solution chemistry unit 8 general chemistry spring ’13
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Objectives (Ch. 15) Understand and describe the basic properties of water
and ice and how they effect the world around you. Explain the high surface tension and low vapor pressure
of water in terms of the structure of the water molecule and hydrogen bonding (15.1.1)
Distinguish between solvent and solute (15.2.1) Describe what happens in the solution process (15.2.2) Explain why all ionic compounds are electrolytes (15.2.3) Distinguish between suspension and solution (15.3.2) Identify the distinguishing characteristic of a colloid
(15.3.2)
Simply Water
Ice has a low density. (Does ice float?)
It’s a polar molecule Slightly positive (+) on one end Slightly negative (-) on another
Look what it does to salt! It also easily bonds to itself and
easily pulls compounds apart
Polar vs. Non-polar?
It’s like a tug of war… Even pulling of
electrons means it’s non-polar
Uneven pulling means it’s polar
Water’s Hydrogen Bond
Water molecules are not connected by full covalent bonds, but they’re pretty strong
The formation between the hydrogen atoms on one molecule and a highly electronegative atom on another is called a hydrogen bond.
Atoms that can do this are hydrogen, oxygen, fluorine, and nitrogen
Keep reading…
Hydrogen Bonds cont.
Any molecule with O-H bonds has the potential to form hydrogen bonds.
Alcohols (molecules with O-H bonds) also form hydrogen bonds. Have similar properties to water
Proteins, nucleic acids, and carbohydrates also can form hydrogen bonds. How they form and shape
determines how they’re used biologically
Water revisited
Boils at 100oC Freezes at 0oC
Expands due to hydrogen bonding Solid state is highly organized
One drop = ~2x1021 molecules
Water in the Solid State
The structure of ice is a regular open framework of water molecules arranged like a honeycomb. Add energy and the
framework collapses The molecules are closer
together making water more dense than ice
Implications on aquatic life?
Surface Tension
Surface tension: the inward force, or pull, that tends to minimize the surface area of a liquid Causes drop to pull together All liquids have a surface tension
Mercury has high surface tension
Surfactants Interfere with H-bonding and reduce
surface tension Soaps and detergents
Capillarity
Results from the competition between the attractive forces between the molecules of the liquid and the attractive forces between the liquid and the tube that contains it
Vapor Pressure of Water
Results from the molecules escaping from the surface of the liquid and entering the vapor phase.
In water, H-bonds hold on to water molecules tightly Tendency for molecules to escape is low Evaporation of water is low What would happen if it was fast?!
Specific Heat of water
Specific heat: measures the amount of heat, in Joules, needed to raise the temperature of 1g of substance by 1oC.
For water it’s 4.18, pretty high. Why it takes so long to boil water It takes a long time to absorb or release more heat
for its temperature to change 1oC than a lot of other substances.
Think of a pool in the summer time.
Water: The Universal Solvent
Almost always found in solution A very good solvent due to its polar abilities Examples of aqueous solutions
Milk Soda pop Coffee and tea Tap water Look at the ingredient list of a liquidy beverage.
Water is probably there.
Solutions
When one substance dissolves into another, that is called a SOLUTION Example: sugar water, Kool-Aid
There are two main parts of a solution: SOLUTE= the dissolved material
Example: sugar, salt, oxygen (air) SOLVENT= the substance that is doing the
dissolving (usually a liquid) Usually present in the highest amount Example: Water, nitrogen (air)
Dissolving
Water, a polar molecule, is capable of dissolving a range of compounds
Many ionic compounds (like NaCl) are soluble in water
When dissolved, ionic solutions are very good conductors of electricity
Electrolytes vs. Nonelectrolytes
Electrolyte A compound that conducts an electric
current when it is in an aqueous solution All ionic compounds are electrolytes
because they dissociate into ions Nonelectrolyte
A compound that does not conduct an electric current in aqueous solution
Many covalents are this because they are not composed of ions
The electric pickle
How does dissolving happen?
Ionic solids are composed of positive and negative ions.
Water has a positive and negative end (it’s polar) Opposites attract and the ionic compounds
separate into ions. The process by which charged particles in an
ionic solid separate from one another is solvation
Dissolving Covalent substances
Sugar is the best example Almost 200 grams can dissolve in 100 mL of H2O! It has O-H bonds, which makes it polar, so it’s easily
dissolvable in water If a molecule contains O-H bonds, it will tend to be
polar and it can form hydrogen bonds with water.
Dissolving Covalent substances
Covalent molecules are simply separated from one another by water molecules.
They don’t solvate into separate ions Covalent solutions can’t conduct electricity
“Like Dissolves Like”
This means that dissolving occurs when similarities exist between the solvent and the solute.
Sugar is a polar molecule, so is water, and water tends to dissolve substances that are polar or that form hydrogen bonds.
Oil and water don’t mix. Oil is nonpolar But different oils are “like” enough to mix and stay
mixed. Oil and gasoline.
Suspensions
A (heterogeneous) mixture from which particles settle out upon standing
A suspension differs from a solution because the particles are much larger and do not stay suspended indefinitely Cornstarch mixed with water thickens
sauces
Suspensions
Two substances are clearly identified Dispersed phase
Ex) clay, dirt, sand, flour, CO2
Dispersion medium Water, ethanol, Milk or Cream
Think of a glass of water with sand or mud in it. Typically easy to separate
Colloids
A heterogeneous mixture containing particles that range in size from 1nm to 1000nm
Particles spread throughout the dispersion medium (s, l, g) Glues Gelatin Paint Aerosol sprays smoke
Tyndall Effect, Coagulation
Tyndall Effect The scattering of visible light by colloidal particles Suspensions also do this but solutions don’t.
Coagulation The clumping of particles in a colloid
Emulsions
A colloidal dispersion of a liquid in a liquid
Must have an “emulsifying agent” To form the emulsion To maintain stability Ex) soap, detergent, egg yolk
Objectives (Ch. 16)
Identify the factors that determine the rate at which a solute dissolves (16.1.1)
Identify the units usually used to express the solubility of a solute (16.1.2)
Identify the factors that determine the mass of solute that will dissolve in a given mass of solute (16.1.3)
Solution Formation
Determining factors Composition of solvent Composition of the solute
Speed of dissolving factors Stirring (agitation) Temperature Surface area
All involve contact between solvent and solute
Rate of Dissolving
Stirring the Solution Increases the interaction
between water molecules and the solute.
Solute and solvent interact more often, the rate of dissolution is faster.
Does not influence the amount of solute that will dissolve
Oil will never mix with water not matter how long you stir or shake that Italian dressing
Rate of Dissolving
Heating the Solution Increases kinetic energy
of the water molecules Increases frequency and
force of the collisions between solute and solvent
Rate of Dissolving
Grinding the solute Creates more surface area
(remember the big fireball demo?)
Solvent molecules attack the edged surfaces of solute crystals.
The more surface area exposed, the faster the rate of dissolving
Solubility
When solute enters the solvent… Particles move from the solid into the
solution Other dissolved particles move from
solution back to the solid Occurs at the same rate Called a saturated solution
Will stay this way as long as temperature stays the same
Solubility
The amount of solute that dissolves in a given amount of solvent at a specified temperature and pressure to produce a saturated solution
Units: grams per liter (g/L) Miscible
two liquids that dissolve in each other Immiscible
Two liquids not soluble in each other
Types of Solutions
Saturated solution Solution holding the max. amt. Of solute
per amt. Of solution under given conditions.
Add more solute it won’t dissolve Unsaturated solution
The amt. of solute is less than the max that could be dissolved.
Add more solute it will dissolve
Solutions
Supersaturated solution Contain more solute than the usual max.
amt. And are unstable. Add a crystal and it fills the container with
crystals
Effects of Pressure Huge effect on gases, very little on solids and
liquids Gas solubility increases as the partial pressure
of the gas above the solution increases. (direct relationship)
Ex) Soda bottle has lots of dissolved CO2 in it which is forced in at the plant.
When you open the bottle you hear a hiss and CO2 starts escaping from the bottle decreasing the concentration of CO2 in the bottle
Solubility Curves
The solubility of substances changes with temperature For example, is it easier to dissolve sugar
in hot or cold coffee? Solids become more soluble at higher
temperatures Gases become less soluble at higher
temperatures
Solubility Curves (cont.)
Scientist have studied many substances solubility at different temperatures They created
graphs which show this data
Solubility Curves (cont.) Let’s simplify the graph with all the
substances down to just one substanceSolubility of KCl in 100 g of water
0
10
20
30
40
50
60
0 10 20 30 40 50 60 70 80 90 100
Temperature (degrees Celcius)
Gra
ms
of s
olut
e di
ssol
ved
in 1
00 g
of
wat
er
Solubility Curves (cont.)
Is KCl a solid or
gas in this graph?
Solubility of KCl in 100 g of water
0
10
20
30
40
50
60
70
0 20 40 60 80 100 120
Temperature (degrees Celcius)
Gra
ms o
f so
lute
dis
so
lved
in
100 g
of
wate
r
Solubility Curves (cont.) How many grams of KCl will dissolve in
500g of water at 80°C? 260g of KCl (52g x 5 = 260g)
Solubility of KCl in 100 g of water
0
10
20
30
40
50
60
70
0 20 40 60 80 100 120
Temperature (degrees Celcius)
Gra
ms o
f so
lute
dis
so
lved
in
100 g
of
wate
r
Solubility Curves (cont.) How many grams of water will it take to dissolve
26 g of KCl at 80°C? 50g of H2O (1/2 of what dissolves in 100g H2O)
(% of 100g: 26g/52g=.50)
Solubility of KCl in 100 g of water
0
10
20
30
40
50
60
70
0 20 40 60 80 100 120
Temperature (degrees Celcius)
Gra
ms o
f so
lute
dis
so
lved
in
100 g
of
wate
r
Solubility Curves (cont.) If one dissolves 95 grams of KCl in 250
grams of water at 80°C, what kind of solution will they have? Unsaturated
Solubility of KCl in 100 g of water
0
10
20
30
40
50
60
70
0 20 40 60 80 100 120
Temperature (degrees Celcius)
Gra
ms o
f so
lute
dis
so
lved
in
100 g
of
wate
r
You need to determine the saturation point before you can decide the type of solution.
(Sat. pt. from graph)x(%H2O)
52 g x 2.5 = 130 g
Saturation point in 250g of water
Practice Problem #1
How many grams of NH4Cl will dissolve in 300 grams of water at 70°C?
If one dissolves 137.5 grams of NaNO3 in 125 grams of water at 45°C, what kind of solution will they have?
186g NH4Cl (62g x 3)
Saturated110g x 1.25= 137.5 The same as what it asks
Objectives
Solve problems involving the Molarity of a solution (16.2.1)
Describe the effect of dilution on the total moles of solute in solution (16.2.2)
Define percent by volume solutions (16.2.3)
Solution Concentration
What does concentration mean? It tells us how much solute is dissolved in a
given volume of solution “dilute solution” has a small amount of
solute “concentrated solution” has a large amount
of solute Both are relative and are very imprecise Qualitative… not quantitative
Molarity
You can describe the precise concentration quantitatively with Molarity.
Molarity (symbol M). Relates the amount of solute to a given volume of
solution. The number of moles of solute dissolved in one
liter of solution. Ex) a solution labeled 3M NaCl is read “three molar
sodium chloride solution”
Molarity Expressed in this manner:
How do you convert from grams to moles?Divide by molar mass foo!
Hint: what is the equation for density?
mass
Density volume
Molarity
Drain cleaner is made with caustic sodium hydroxide, NaOH. The Dow company prepares a bottle of drain cleaner from 24.0 g of NaOH dissolved in 0.100 L of solution. What is the molarity?
Molar mass of NaOH (40.00 g/mol)
24.0g NaOH 1 mol NaOH
mol NaOH
= L solution = 6.00M NaOH0.100 L solution 40.00g NaOH
We want Liters, leave this alone Units for molarity
Molarity- Writing Unit Factors
We can solve molarity calculations by using the solution concentration as a unit factor.
Example: 6.00M solution of NaOH contains 6.00 mol of solute in each liter of solution. Written as:
6.00 mol NaOH or 1 L solution
1 L solution 6.00 mol NaOH
There’s 1000mL in 1L
Preparing Solutions
How would you prepare 1.0L of a 0.15M sodium chloride solution? I.O.W. How many grams NaCl are needed?
Think: First, determine the mass of NaCl to add to a 1.0 L container. The 0.15M solution must contain 0.15 moles of NaCl per liter of solution (definition of molarity).
1. Use molarity to convert to moles.
2. Then use molar mass to go from moles to grams.
1.0 L solution 0.15 mol NaCl 58.5 g NaCl = 8.8 g NaCl 1 1 L solution 1 mol NaCl
Making that solution
1. Obtain a volumetric flask
2. Measure 8.8 g of NaCl
3. Add solute to a small amount of water, about 300 mL, to dissolve
4. Add enough additional water to bring the total volume to 1.0 L, to the etched line on flask
Preparing a Different Volume of Solution
How would you prepare 5.0 L of a 1.5M solution of glucose, C6H12O6
Think: You need to determine the number of grams of glucose to add to a 5-L container. The 1.5M solution has 1.5 mol of glucose (use this to convert to grams.
5.0 L solution 1.5 mol glucose 180 g glucose
= 1400 g glucose 1 1 L solution 1 mol glucose
MOLARITY MOLAR MASS
Making Dilutions
Stock solutions of acids are very concentrated HNO3 comes in 15.8M
H2SO4 comes in 18.0M HCl comes in 12M
This does not tell how “nasty” these are…that’s a different unit
Using an acid in this concentration is incredibly dangerous
I dilute them for lab purposes… how?
Dilutions
Diluting a solution reduces the # of moles of solute per unit volume The total number of moles of solute in solution does
not change mol solute before dilution = mol solute after dilution
Equation for dilutions: M1 x V1 = M2 x V2
M1 and V1 are initial readings M2 x V2 are for the diluted solution
Units of volume must match (mL or L)
Dilution example How many milliliters of 2.00M MgSO4 must be diluted with
water to prepare 100.0mL of 0.400M MgSO4? M1= 2.00M M2=0.400M
V1= ? V2= 100mL M1 x V1 = M2 x V2
Dilution answer
Substitute and solve for V1
20 mL Take 20mL of initial solution and dilute with
enough water to increase volume to 100 mL
DO NOT ADD 100mL of water, this will give 120mL of solution, not 100mL
Dilution Example #2
If you take 10mL of a 3.42M solution and dilute to 100mL, what is your new concentration? M1 x V1 = M2 x V2
(10mL)(3.42) = (M2)(100mL)
M2 = 0.342M
Dilution Example #3
You need a 2.10M solution. You need 1500mL of it. How much of a 12M solution do you need to use? M1 x V1 = M2 x V2
(12M)(V1) = (2.10M)(1500mL)
V1 = 262.5mL
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